Chemical Bonding and Molecular Structure · Lesson 5 of 12
Bond Polarity and Molecular Dipole Moments
“Use electron attraction and molecular geometry to explain polar bonds, net dipoles, and the overlap between ionic and covalent character.”
• Distinguish nonpolar and polar covalent bonds using electronegativity. • Calculate a dipole moment from charge and separation and interpret its units. • Explain why bond dipoles may cancel in a molecule and compare ammonia with nitrogen trifluoride. • Use polarisation and Fajans’ rules to compare covalent character in ionic compounds.
Sharing does not always mean sharing equally
A covalent bond contains shared electrons, but the atoms do not always attract them equally. In a bond between identical atoms, the two nuclei have the same attraction for the shared density. In a bond between different atoms, the electron density can be displaced towards one atom. This unequal distribution gives the bond polarity.
The tendency of an atom in a bonded chemical environment to attract shared electron density towards itself.
In H₂, Cl₂, and F₂, the two identical atoms do not create a permanent bond dipole. In HF, fluorine attracts the shared density more strongly than hydrogen. Hydrogen becomes partially positive and fluorine partially negative. These are partial charges, written with delta symbols, rather than the full charges of separated ions.
Electronegativity describes an atom in a bond. Electron-gain enthalpy describes an energy change when an isolated gaseous atom accepts an electron. The ideas can be related, but they are not the same quantity: one is a tendency used to interpret bonds, the other a specified gaseous process with energy units.
A covalent bond with an unequal distribution of electron density, producing partial positive and negative ends.
A polar covalent bond is not a pair of completely separated ions. Ionic compounds can also show shared electron density. The labels “ionic” and “covalent” identify useful limiting descriptions, while many real bonds have mixed character.
Measuring charge separation
A dipole consists of opposite charges separated in space. Its dipole moment depends both on the amount of charge and on the distance between the charge centres. A small charge separation over a large distance can therefore have the same moment as a larger separation over a shorter distance.
A dipole moment has a direction, so it is a vector. The conventional electric dipole vector points from negative charge towards positive charge. Chemistry often uses a crossed bond arrow in the opposite direction: the crossed tail indicates the positive end and the arrowhead points towards the electron-rich negative end. Always identify the convention before reading a diagram.
Problem
A bond has charge separation over . Calculate its moment in SI units and debye.
- 1.Multiply charge by distance: .
- 2.Convert to debye by dividing by the value of one debye: .
- 3.The result describes the magnitude of the moment. Its direction requires knowing which end is electron rich.
Problem
Water has dipole moment 1.85 D. Express it in coulomb metres.
- 1.One debye equals .
- 2.Multiply: .
- 3.To three significant figures, . This is the net moment of the whole molecule, not the magnitude of one isolated O–H bond dipole.
Molecular shape decides how bond dipoles combine
For a molecule with several bonds, adding the magnitudes of the individual bond dipoles usually gives the wrong result. Their directions matter. The net molecular moment is a vector sum of the charge distribution, often interpreted through the individual bond dipoles and the effects of lone-pair electron density.
In linear BeF₂, the two identical Be–F bond dipoles point in opposite directions and cancel. In trigonal planar BF₃, the three equal dipoles are separated by 120°; their vector sum is zero. Symmetrical tetrahedral CCl₄ also has zero net dipole. In each case the bonds can be polar even though the molecule is nonpolar.
In bent water, the two O–H dipoles do not oppose each other along a straight line. Their components along the molecular bisector reinforce, producing a nonzero result. This is why the molecular geometry is essential when interpreting polarity. Here these shapes are given; the next lesson explains how to predict them.
Problem
Why can carbon dioxide have zero dipole moment although its C–O bonds are polar?
- 1.Oxygen attracts bond electron density more strongly than carbon, so each C–O bond has a dipole.
- 2.The molecule is linear and the two bonds are equivalent. Their dipoles have equal magnitudes and opposite directions.
- 3.The vector sum is zero. Zero net molecular dipole therefore does not mean that every bond is nonpolar.
Why ammonia has a larger moment than nitrogen trifluoride
NH₃ and NF₃ both have pyramidal shapes with a lone pair on nitrogen. It may seem that the strongly electronegative fluorine should make NF₃ the more polar molecule. That conclusion ignores the directions of the contributions to the molecular moment.
In the chapter’s bond-dipole picture of NH₃, the resultant of the N–H bond dipoles acts in the same general direction as the lone-pair contribution. They reinforce. In NF₃, the N–F bond-dipole resultant acts against that contribution. Partial cancellation gives NF₃ a much smaller net moment. The comparison illustrates why electronegativity difference alone cannot rank whole-molecule dipoles.
| Molecule | Shape | Representative dipole moment (D) |
|---|---|---|
| HF | Linear diatomic | 1.78 |
| HCl | Linear diatomic | 1.07 |
| HBr | Linear diatomic | 0.79 |
| HI | Linear diatomic | 0.38 |
| H₂ | Linear diatomic | 0 |
| H₂O | Bent | 1.85 |
| H₂S | Bent | 0.95 |
| CO₂ | Linear | 0 |
| NH₃ | Trigonal pyramidal | 1.47 |
| NF₃ | Trigonal pyramidal | 0.23 |
| BF₃ | Trigonal planar | 0 |
| CH₄ | Tetrahedral | 0 |
| CHCl₃ | Tetrahedral around carbon | 1.04 |
| CCl₄ | Tetrahedral | 0 |
Polarisation gives ionic bonds covalent character
Now start from the ionic limit. A cation attracts the electron cloud of an anion and can distort it towards itself. This distortion is polarisation. Electron density drawn into the region between the nuclei increases the covalent character of the interaction.
Polarising power is a cation’s ability to distort another ion’s electron cloud. Polarisability is the ease with which an anion’s cloud can be distorted.
A small cation concentrates positive charge in a small region and generally polarises strongly. A large anion has a more easily distorted outer cloud. Greater cation charge also increases polarising power. Fajans’ rules combine these trends to compare covalent character among otherwise related ionic compounds.
| Change | Expected effect on covalent character |
|---|---|
| Smaller cation, otherwise comparable | Usually increases |
| Larger anion, otherwise comparable | Usually increases |
| Higher positive charge on cation | Usually increases |
| Comparable cations with a filled d-electron arrangement rather than a noble-gas outer arrangement | Often greater polarisation and covalent character |
The electronic-configuration qualification matters because ions with similar size and charge need not shield nuclear attraction equally. The chapter contrasts typical transition-metal-type filled-d arrangements with noble-gas configurations. Treat this as a comparison under similar conditions, rather than a substitute for considering size and charge.
Problem
Compare the expected covalent character of LiF and LiI using Fajans’ rules.
- 1.The cation is the same, , so the main comparison is between the anions.
- 2.Iodide is larger and more polarisable than fluoride. Its electron cloud is more easily distorted by the cation.
- 3.LiI is therefore expected to show greater covalent character than LiF. This is a qualitative comparison, not a calculation of an exact percentage.
Check your understanding
First distinguish a bond from the whole molecule, then decide whether geometry or electron-cloud distortion is the relevant idea. This prevents several common polarity errors.
Quiz
Which molecule has polar bonds but zero net dipole in its stated symmetrical shape?
If Q is unchanged and r doubles, the dipole moment:
Why is NF₃ less polar overall than NH₃?
Which combination most favours anion polarisation?
Electronegativity differs from electron-gain enthalpy because it:
Practice Problems
- Use HF and H₂ to explain polar and nonpolar covalent bonds.
- Calculate the dipole moment for and , first in coulomb metres and then in debye.
- Sketch bond-dipole directions for H₂O, BeF₂, BF₃, and CCl₄ and explain which sums cancel.
- Explain why NH₃ and NF₃ cannot be ranked by considering only the electronegativity of fluorine.
- Compare LiCl and NaCl qualitatively using cation size, and LiF and LiI using anion size.
- Distinguish formal charge, partial charge, and the net charge of a molecular ion.
- Explain how a measured nonzero dipole can help rule out a proposed symmetrical shape with identical surrounding atoms.
Key Takeaways
• Electronegativity differences can make shared electron density unequal. • Dipole moment depends on charge separation and is a vector quantity. • A molecule can have polar bonds yet a zero net dipole because of symmetry. • The NH₃/NF₃ comparison requires the directions of all relevant contributions. • Fajans’ rules relate cation polarising power and anion polarisability to covalent character.